Method of determining ph and buffers and indicators therefor



S. F. ACREE Original Filed May 2, 1929 6 Sheets-Sheet l w Mw METHOD 0FDETERMINING pH AND BUFFERS AND INDICATORS THEREFOR Nov'. 7, 1939. s FACREE N 2,178,550

METHOD 0F DETERMI'NING pH @ND BUFFERS AND INDICATORS THEREFOR origina-1Filed may 2, 1929 e sheets-'sheet 2 Neyo 7 1939., s. F. AGREE METHOD 0FDETERMINING pH AND BUFFERS AND INDICATORS THEREFOR Qgginal Filed May 2,1929 6 Sheets-Sheet 5 A TTORNEYS.

Nov. 7, 1939. s. FL ACREE METHOD 0F DETEBMINING pH AND BUFFERS ANDINDICATORS THEREFOR 6 Sheets-Sheet 4 Original Filed May 2, 1929 Wsw .w@Xn INVENTOR.

TTORNEYS.

Nov. l?, 1939. s F, AGREE 2,178,550

METHOD 0F DETERMINING pH AND BUFFERS AND INDICATORS THEREFOR s. F.ACREE" y METHOD 'OFDETERMINING pn AND' BUFFERS AND INDICATORS THERDFOR lOriginal Filed May 2; 1929 6 Slxets--Shet 6 v L. Q m @Q N xo In q 'kb Nv Wam/mf aw# 27AM/v /u/m 4m/0f@ Ms'- z mamas/075.7 -a INVENTOR may y J/J Mw v a Mana /va/ raz/J H j BY lf j A41@ W ATTORNEYS.

raras METHOD F DETERMINING pH AND BUF- FERS AND INDICATORS THEREFORSolomon F. Aeree, Washington, D. C.

Application May 2, 1929, Serial No. 359,930 Renewed January I4, 1939 10Claims. (Cl. 23--230) The present invention comprises a stable hyandshow these acids to meet the conditions neardrogen ion standard (i. e.,H-ion) known as sodily exactly. um acid ortho phthalate, which is at thesame It is therefore clear that phthalic, malic, suctime an excellentbuffer; and it comprises comcinic, glutaric, tartaric, and citric acidsgive acid est " positions thereof with indicators adjusted in hysaltswhose pI-I value in M/ 10 solutions are fixed 5 drogen ionconcentration, and with other buffor each but increase with dilution anddiffer fers. It comprises especially solutions thereof among themselvesand iheeOTe OIIII e Series Of and pI-I comparisons therewith made withsuperpH standards.

pure water of practically pH 7.0 such as I prepare I have found sodiumacid phthalate one of the lo free from Carbon dioxide and traces 0f buerbest of these standards for the following reasons. 1U

compounds by my process in a special type of It can be made cheaply,easily, and Dure from still. soda and orthophthalic anhydride orphthalic In the developments and use of the improved acid. Itcrystallizes with one molecule of water types of hydrogen-calomelelectrode systems deat times et W temperatures elOU-Ild 0 C. but

` l5 scribed in another application Ser. No. 624,277 it `precipitates atordinary temperatures up to boil- `15 was necessary to employ a newhydrogen ion ing with 0.5 molecule of water to form a very stastandardin one hydrogen electrode chamber and ble eeid salt WhiCh keepsindelitely, d0es not measure its electromotive force (hereinafter add orlose water in the air, but is dehydrated el? called E. M. F.) againstthe reference calomel ll0 C. WiihOut further deCOmPOSiUiOll 0T fermefggelectrodes and against the hydrogen electrodes in tion 0f 2m anhydrideof higher DH Value- The 20 the unknown solutions. In other words, Idesired enhydleus Salt keeps dry indefinitely and has to discover andeliminate some of the disadvanthe saine pH as the hydrated sell? irl M/10 solutages of the hydrogen ion standards used theretiens, namely about3.95- The acid salt Can be toore and replace them with one 0r moreStandsterilized without decomposition in either the e,- ardssubstantially free from faults, dry state or in solution by autoclavingat about In searching for a hydrogen ion Standard I 110 C. Its lvl/10solution sterilized alone or with busied myself primarily with basicsalts of organ- M/ 200 hymOl keeps indefinitelyic poly-bases, such asNH2CH2CH2NH2I-IC1, name- The acid salt can be made readily in either ofly, ethylene diamine hydrochloride, and acid tWO Ways. The 10h/Cheli@anhydride een be boiled salts of organic poly-acids such as With waterto form phthalic acid and treated 30 with one equivalent of pure sodiumbicarbonate NaOOCCGH4COOH or sodium hydroxide. The pH in 1v1/1o somtionsodium acid ortho phthalate. I chose especially should be 3.95. When theresulting solution is those acids or bases having ratios of about 10:1evaporated to crystallization and cooled the acid to 100:1 between theseparate ionization consalt crystallizes out and more can be obtained by35 stants so as to get smooth practically flat hydroevaporation of themother liquors. Another gen ion titration curves, without largeinflections, method is to boil the anhydride with about 4-5 forexpressing the relation between the degree of parts of water andneutralize the resulting neutralization of the acid with a base and thephthalic acid with exactly two molecules of sodilcglQVH value which isoften called the pH of the um bicarbonate, or until a test drop of theboiled 40 solution. VH is the number of liters of solution solutiongives a faint pink color with phenol containing one gram equivalent ofhydrogen ions. phthalein. Exactly the same Weight of phthalic With suchsmooth titration curves accidental anhydride is then dissolved in thisboiling soluinequaiities cr impurities of acid or basic charaction ofdisodium phthalate and the solution is ter cannot make such variation inthe pH values filtered hot from impurities when necessary and obtainedduring the preparation and use of such concentrated until the crystalsof sodium acid partly neutralized salts as I-I-ion standards. phthalatebegin to precipitate. The solution is As no exact pH or I-I-iontitration curves were then cooled, stirred and iiltered. The crystalsthen available I rst used the ionization conare then eil" dlied and arepllie but may be re- ;g stents given by conductivities or other methodscrystallized if desired The mother liqllOIS 0n 50 for a number of suchpolybasic organic acids and evaporation give more of the pure saltwithout selected phthalic, malic, citric, succinic, malonic, any furtherchemical pI-I adjustment. tartaric and oxalic acids as the mostpromising. The hydrated or dehydrated salt may be used The actual pHtitration curves which I later to make up standard solutions in mysuperpure 1 3 made are, except for oXalic acid,l very smoothCOz-salt-free Water or may be used to prepare 55 Cil l li Y w' wstandard mixed solid buiers. For example, l cc. portions o1" thesolution M/ l0 (or M/ 1000) sodium acid phthalate sterilized with M/ 200thymol (or with formalin, chloramines, etc.) may be treated with anydesired volume of N/ (or N/lOO) hydrochloric acid or sodium hydroxide upto one equivalent or more and diluted to 20 cc. and checkedelectrometrically to give a series of pH standards of any values (e. g.,0.2 pH apart) between pH 2.0 and pH 7.0 as shown in Fig. 5 of myapplication Ser. No. 624,277 (Fig 2 of the present case).

The sulphonphthalein indicators are the best series yet developedbecause they are soluble enough in we; er to forni for example lvl/1009solutions, other indicators such as phenolphthalein requiring forsolution ethyl alcohol which interferes with the proper color reactionsin pH work. In such aqueous solutions of, e. g. 0.04% concentration, thecolor is very intense red, blue or green, when two molecular equivalentsof alkali are added, the first forming a yellow quinoidal salt and thesecond molecule of alkali forming the intensely colored quinonephenolate salt, according to the following equation:

The color used as an indication of pH values is in this case due to amixture of two diierent color-bearing compounds. The intensely coloredsalts are not hydrated and decolorized by a small excess of alkali suchas 3 to 5 moles used to con-- vert the indicator completely into theintensely colored sait in testing its purity by a spectrophotometer. Bymaking desirable chlor, brom, iodo, nitro, amino, methyl and otherderivatives of the phenolic groups particularly, it is possible to varythe relative ionization constants of the phenolic and sulphonic groupsand to vary the pH range in which the color changes take place, such as11H2 to 10 approximately. The purity of the indicator can be tested bytitration with alkali to see what fraction of one molecule of alkali isnecessary to change the yellow to the red color, for example, withphenolsulphonphthalein. Heretoore these monobasic indicator salts havebeen used as the standard indicator solutions (e. g., 0.04 per cent),without any adjustment of the pH as described herein for isohydrieindicator tests. For example, 6 mgm. or the monosodium salt per cc. wasused for injections into humans and the urine was collected and analiquot portion made alkaline and compared against standardphenolsulphonphthaiein in excess alkali to study` the amount ofphenolsulphonphthalein excreted as a measure or renal efciency. Beforemy work the pH value of this solution was not adjusted to about 7.5 orclose to the pH value of the blood stream or muscle tissues, that is,isohydric therewith. This pil 7.5 is about the midpoint adjustment inthe useful range of this indicator. This adjustment should be done andcan be by careful colorinietric comparisons While titrating theindicator solution with standard acid or alkali or can be doneelectrometrically. These adjusted solutions of indicator salts can beevaporated or the same composition can be made as described with otherbuffers by mixing the ratio of monobasi-c and dibasic indicator saltscalculated from the pH titration curve of the indicator as described fornonchromophoric buffers. When dissolved in pure water these indicatorsalts give the correct pH calculated from and or their adjustment. Theseadjusted solid indicator salts can be mixed in standardized amounts withpure sodium chloride and molded in standard size indicator tablets todissolve in, e. g., l0 or 20 cc. of pure water to give standardindicator solutions. Or these adjusted indicator salts can be mixed instandard amounts with buiers giving the same pH and molded in tablet orother form as standard bufferindicator tablets to dissolve in e. g., l0or 20 cc. of pure water to give standard buffer-indicator solutions ofdeiinite pl-I and color content. I nd it very advantageous to preparethe indicator-salt mass separately, as outlined above for buffers, andhaving a higher concentration of indicator to correspond with the laterdilution with the buffer mass. The builer mass is then prepared for agiven pi-I and mixed with the indicator-salt mass, and part of themixture is pel.n leted. The true pH value of the finished mass isdetermined carefully electronietrically and colorimetrically and anysmall discrepancies in either indicator content or in pH value areadjusted by adding more indicator-salt or buier mixture of the same orother necessary pH value as described under buier tablets. When theentire indicator buier mass is adjusted, it is pelleted and thesetablets then contain the amount of indicator necessary to give the samecolor intensity as I obtain from a standard indicator tablet in asolution of the same pH value. It is therefore very convenient to havesets of these buffer indicator tablets varying about 0.2 pH units apartand covering any desired ranges such as those given by hexamethylpararosaniline chloride (pH -l) thymolsulphonphthalein (pI-I 1.2-2.8),tetrabromphenolsulphonphthalein (pl-l 3.0-4.5), benzoylauramine (pH4.5-5.5), o-carboxybenzeneazodimethylaniline (pH ifi-6.0),o-carboxybenzeneazodipropylaniline (pH 4.8-6.4)

dibromoorthocresolsulphonphthalein (pH 5 .2- 6.8)dibrointhyinolsulphonphthalein (pI-l 6.0- 7.6) phenolsulphonphthalein(pH 6.8-8.4)

orthccresol sulphonphthalein (pI-I 7.2-8.8), alpha naphtholphthalein(pl-I 7.2-8.7) thymolsulphonphthalein (pI-I 8.0-9.6) ortho cresolphthalein (pH 8.2-9.8), para-nitro-benzene-azosalicyclic acid (pH10G-12.0), triphenylrosaniline sulphonic acid (pH 11-13) and alizarineblue sodiuln bisulphite (pH .l2-14). It is possible that the salts ororganic or other materials in industrial or natural liquids, urine,biological fluids, special soil Waters, etc., etc., may have aparticular influence on the colors of the specific indicators and makethe colorimetric readings seem different from the electrometric pHvalues. These differences are called salt or protein. errors. It is insuch cases possible by my methods to adjust the indicator content and pHvalue or the standard indicator tablets and buffer-indicator tablets toovercome these discrepancies and give useful tablets which will givetrue pH values independent of the specific inuence of the particulardisturbing compound or compounds in the solution under test. .All ofthese indicator tablets or buffer tablets or buffer-indicator tablets orampoules of the indicators standardized amounts and pH values, such asthe phenolsulpl'ionphthalein at 19H75 for renal tests, can be sterilizedby heating in vials at 110 C. on three successive days for use insterile water, culture media, or other solutions whose ph is to betested and/or regulated.

My IVI/1000 or IVI/2000 sulphonphthalein indicators of the same pI-Ivalues diluted therewith to about M/50,000 indicator strength make eX-cellent pI-I-indicator standards free from salt errors. For example, Ican adjust these sulphonphthaleins, or other similar indicators, to anydesired series of pH values in the range of each indicator by dissolvingthem in pH 'l CO2-saltfree superpure water to make M/ 1000 or M/2000solutions and titrating portions of them (e. g., cc.) with standard N/10 acid or alkali while making careful colorimetric comparisons thereofagainst a set of companion micro buiIer-color standards of the samedesired pH series; or the adjustment can be done electrometrically witha hydrogen electrode or quinhydrone electrode. I can thus make a .seriesof standard M/1000 or M/2000 solutions of any indicator 0.2 pI-I apartand covering the'useiul pI-I range of that indicator. In the claims andelsewhere, unless otherwise specifically noted, the word adjusted andrelated nouns and verbs are intended to mean that the colorless buiTersand the colored or chromophoric (colo-r bearing) buffers calledindicators are treated in known amounts with known amounts of acid oralkali and thereby adjusted in pH value by this described electrometricor colorimetric procedure to give a desired or predetermined pH value,which in certain cases is also freed from any errors due to salt errorsor salt or protein eiects. These indicators are chroinophoric orcolor-bearing buffers Whereas the usual phosphate, borate, phthalate andother colorless buiers are called nonchromophoric or non-colored ornon-color bearing buffers.

By diluting successive small volumes of each oi these standard adjustedindicator solutions (e. g., 0.2 cc.) with 1/5 to 1 volume oi anisohydric buffer solution (i. e., of the same pI-I value) Whileadjusting the indicator pH values, a comparison series of accuratebuffer-color micro standards of any desired buffer concentration (NI/10to IVI/1000) is made for checking 0.2 cc. portions of the correspondingM/ 1000 or lli/2000 indicator solutions themselves during adjustment asabove; both solutions are viewed in suiciently thin layers in small 1cc. Pyrex dishes, or stoppered Pyrex test tubes laid on their sides, andthrough the same cross section in accordance with Beers law. These microbuffer-color standards are also used when testing small quantities ofunknowns with e. g., 0.2 cc. portions oi the adjusted series or the sameindicator.

In the adjustment of the pI-I values of the standard sulphonphthaleinindicator solutions themselves it is convenient to dissolve 0.1 gram inone equivalent oi alkali to neutralize the Sulphonic acid group (e. g.,with a standard alkali such as lvl/20 sodium hydroxide) and then dilutethis solution to 0.02 or 0.04 percent or 1v1/1000 or IVI/2000 or anyother concentration desired. Such a solution has a pH value at or nearthe lowest useful pI-I of the indicator, with the exception of very acidindicators like tetrabrornphenolsulphonphthalein (brom phenol blue)which requires only a half equivalent of alkali. Amphoteric indicatorslike the hydrochloride or methyl red(para-dimethylamino-azo-benzeneortho-carboxylic acid) require onemolecular equivalent of alkali to neutralize the HC1 group and giveapproximately the lowest useful pI-I value. Portions (e. g., 100 cc.) ofthese standard indicator solutions are then adjusted in pH Values by thecareful addition of further volumes of N/20 sodium hydroxide andchecking against the micro buffer-color standards described herein. Itmust be understood that commercial indicators vary in purity and thatsomewhat variable quantities of N/20 alkali will be required to give thesame pI-I value with diierent sample of the same indicator. Thefollowing figures may be used to illustrate this principle with bromthymol blue, whose adjusted solutions I use to test the purity or pI-Ivalue of the superpure pI-I '7 .0 water made from my still described inthis application and the pI-I values of weakly buffered solutions madefrom such water. I shall give the pI-I value sought, the theoreticalnumber of cc. of N/20 alkali required per 0.1 gram of the indicator, andthe number of cc. found for diierent samples of the indicator. For pH6.2, theory 3.61 cc.; found 3.45, 3.61, 4.10 cc. For pI-I 6.4, theory3.84 cc.; found, 3.70, 3.84 cc. For pH 6.5, theory 4.13 cc.; found,3.95, 4.13, 4.30 cc. For pI-I 6.8, theory 4.44 cc.; found, 4.20, 4.44,4.80 cc. For pI-I 7.0, theory 4.80 cc.; found, 4.65, 4.80, 5.00 cc. ForpH 7.2, theory 5.21 cc.; found, 4.95, 5.21, 5.15 cc. For pI-l theory5.53 cc.; found, 5.45, 5.53 cc. For pI-I 7.6, theory 5.79 cc.; found,5.52, 5.79 cc. I have likewise found the amounts of alkali needed foreach pH value of brom phenol blue, brom cresol green, brom cresolpurple, brom thymol blue, cresol red, phenol red, methyl red, andcorresponding derivatives of tetrachlorand tetrabrom-sulphonphthaleins,and thereby prepared adjusted indicators from pI-I 3.4 to 9.0 in 0.2 pHsteps. The above disclosed method enables anyone skilled in this art toadjust the pH of his own indicator solutions.

By diluting 0.2 cc. of each adjusted indicator solution 1:49 with anydesired buffer such as M/10 to M/1000 sodium acid phthalate or otherbui-Ier adjusted to the same pH as the indicator, I obtained a series ofaccurate buffer-color testtube or 10 cc. standards for making acolorimetric comparison and measurement of the correct pI-I value of anyother buffered or unknown solution available in larger volume and alsodiluted 49:1 with the indicator of the same pI-I. This correct pH ofeven very dilute and weakly buered unknown solutions including mysuperpure conductivity water will generally be obtained by one or twopreliminary tests with the indicator adjusted at its lowest, highest,and mid-pH points and one iinal check on a separate sample of theunknown with the indicator having the pH given by the preliminary tests.this method my isohydric indicator method of using the adjusted pHseries of any indicator together with the adjusted companion series oibuer-color standards. These indicator, butter, and buiTer-colorstandards are best kept sterile with IVI/200 thymol, formaldehyde, etc.,preferably admixed therewith beforehand and included definitely in thepI-I adjustment.

These indicator solutions or any buffer solutions can be protectedagainst changes by the carbondioxide (CO2) of the air by the simple eX-pedient of adding sufficient sodium carbonate and/or bicarbonatecalculated to form in solution the concentration of CO2 required for airequilibrium. The mass law equation CO2 X 3 X 10-7/H2NaI-ICO3 1.2 X 10"5X 3 X 10-7/H=NaHCO3 gives the concentration of sodium bicarbonate I call(NaHCO3) in mols per liter needed for each concentration of hydrogenions (H) to keep these solutions in equilibrium with the atmospheric CO2when 3 10'l is the ionization or activity constant of carbonio acid inthe particular solutions. As discussed on page 6, column l, line 44,through page 7, column 1, line l0, oi this specification, thisstabilization of solutions containing buiier materials against changesby atmospheric carbon dioxide can also result from the establishment ofthe proper carbonate-bicarbonate-carbon dioxide concentrations by simpleaeration.

As illustrated in Fig. '7, curves A, B and C I have found that the usualM/lO to M/20 phthalate, phosphate, borate, etc., buiers give colors withsulphonphthalein indicators which correspond to pH values about 0.2-0.3pH higher (but lower with methyl red in some cases) than the E. M. F.pI-I value, which abnormal colorimetic salt eiiect or salt errordisappears with dilution to IVI/1000 or thereabouts, I have madebuffercolor micro and test tube standards of M/ 1000 concentration orthereabouts which are free of salt errors for my tests against extremelydilute solutions and water, or I have madetheproper corrections for salterrors while adjusting the pH of more concentrated lvl/l0 to M/lOObutler-color standards for use in testing unknown solutions of M/ 10 toM/ 10,000 concentrations.

The anhydrous sodium acid phthalate may be mixed with the properquantities or phthalic acid and anhydrous disodium phthalate, as shownin the titration curve for phthalic acid, Fig. 5 in my application Ser.No. 624,277 (Fig. 2 of the present case), and ground and intimatelyadmixed in a ball mill to give any desired mix for making small tabletsin a pelleting machine or for use in larger quantities in industrialprocesses. By grinding about equal weights of these mixed phthalate orother buffers and sodium chloride in a ball mill the buier tablets canbe easily pelleted in any desired size such as 0.4 gram for solution in20 cc. of water as a pH standard. It is easy to adjust these phthalatemixes to make buiTer tablets for any desired pH in the range of phthalicacid, and hence to make a series of these buffer tab-lets 0.2 pH apart.Likewise these phthalate salts can be mixed with phosphate and beratesalts in the proper proportions to makeuniversal solid and liquid buffermixes by the methods, and Figures 4, 5 and 6 given in my applicationSer. No. 624,277 (Figs. 1, 2 and 3 of the present case).

For example, the sodium acid phthalate may be mixed with thesepredetermined quantities of solid phosphates, thymol, borates,pyrophosphates, phenolsulphonates, etc., or their solutions to giveuniversal buffer solutions or universal buffer tablets orindustrial-size mixtures, as set forth in the above applica-tion. Properquantities of these materials with ionization constants decreasing inthe ratio of about 1:10 to 12100 give a smooth titration curveillustrated in Fig. 6 in Ser. No. 624,277 (Fig. 3 of the present case).By addition of the indicated volumes of N/5 alkali or acid to 10 cc. ofsuch a sterile universal buil-er solution and dilution to 20 cc., xed pHand pI-I indicator standards can be prepared very quickly from such achart. Likewise solid buffer tablets or mixes of these chemicalscovering wide pH values can be made, e. g., pH 2 to l0. These universalbuffer solutions and tablets can be premixed with accurate quantities ofindicators already adjusted to the same pI-I,l such as thesulphonphthalein salts, to give buffer-indicator mixtures suitable tothe pH range of each indicator.

The theory of pH titration curves rests on the fact that when an acid istitrated with a base, the

hydrogen ion concentration CH and the pH Value (pHz-logmCH) graduallychange, the former decreasing and the latter increasing. The CH or pHcan be measured by means of a hydrogen electrode or the standardindicators herein def scribed. The relation between the fractionalequivalents of alkali and the total hydrogen lon concentration (e. g.,CH, H-ion, or H+ or Ht) and the ionization constant or constants of themonoor polyvalent acid can be plotted in graphs and curves and expressedin equations. For example, the separation ionization constants of apolybasic acid like phthalic or citric acid can'be measured by the useof the equations and tion is and KOI-lt the total equivalent of alkaliadded.

in plotting the data ior these equations in graphs and curves it is lessconvenient to plot I-It than logHt or -logI-It (i. e., pH). Figures 1, 2and 3 represent three dierent ways in which the relation between pL-Iand other data may be plotted. Fig. l shows the rise in pH ordinateswhen a monobasic acid with ionization constant K=10-7 is titrated withincreasing fractions of an equivalent of alkali plotted as abscissaefrom right to left. when l0 cc. of a universal buffer solution describedherein are treated with the volumes of standard acid or alkali given asabscissae to right or left of the central zero mark. Figure 2 gives asordinates the pH data for phthalic acid titrated with 2.0 equivalents ofalkali plotted as abscissae from left to right. Curve AB gives pH versusequivalents of total acid salt, KHAnt. Curve A is the acid salt curvethat would be obtained if the iirst acid group could be titrated beforeany neutralization of the second acid group. Curve C is the dibasic saltcurve that would be obtained if the second acid group could be 'tratedby itself. Curve AC is the pH curve actually obtained because the secondacid group is titrated along with the irst acid group after about 75 percent of the latter is neutralized. As indicators are merelycolor-bearing or chromophoric acids or bases, and hence form buffersalts, the pH curves such as Figs. l, 2 and 3 apply equally well toindicators. It is noted that the ends of the pI-i curves for each acidgroup form sharp inflections and that only the middle 50 per cent ispractically straight.

The theory and practice involving such buffer mixtures rest on the factthat during the neutralization of the last 25% of a given acid with abase the pH does not rise sharply as shown in curve A, Fig. 2, becausethe rst part such as 25% of the neutralization of the next weaker acidor acids Figure 3 shows the change in pH Will also be taking place andWhose pH will therefore be kept down as illustrated in curve C, Fig. 2.The actual pI-I titration curve will therefore be a resultant of the twoor more individual pH titration curves as shown in Fig. 5 in Serial No.624,277 and curve AC in Fig. 2 of the present case. Since the middle 5Gper cent portion of the pH titration curve for each acid (or base) ispractically a straight line the difficult part of the calculations andchoice of buier acids or bases is to use the mixtures whose overlappingends of the pI-I curves will also give resultant practically straightlines forming continuations of the middle straight portions of thecurves and hence practically a straight line through the useful part ofthe total curve as shown in Fig. 3. The choice of a mixture lof bufferacids With suitably varying ionization constants should give a smooth pHtitration curve throughout the entire neutralization of such mixture ofbuffer acids with bases, or similarly of buier bases (amines,aminoacids, etc.) with acids; in other words there would be a gradualvariation in pH, making such a single solution an excellent pH standardcovering a wide pH range. A proper choice of such buffer mixtures can beso made, particularly by using iractions of molecular or equivalentquantities of the buffer acids or bases, that the resultant titrationcurve is a smooth or straight line titration curve. The addition ofdisinfectants such as thymol (preferably N /200) or para sulphonic acidcarbolater (p-HOCHiSOaH), preferably M/ 1l), may be used also to formpart of the titration curve andalso to keep the solution sterile toprevent changes in pH of the solution through bacterial action.Furthermore, neutral preservations vsuch as 1-5% formaldehyde,chloramine derivatives, etc., may be used with success but boric acidsolutions did not prove as permanent as was desired. This inventionembraces the preparation of sterile, indefinitely stable universalbuffer solutions of any desired pH values, from which any chosen numberof standard solutions of denite pH values can then be prepared by theaddition of standard acid or alkali, with or without water.

' It has proved desirable to use titration curves as set forth above,one of which is illustrated in Fig. 6, Serial No. 624,277, and Fig. 3 ofthis application, and is formed by plotting as abscissae the volumes ofstandard (e. g., N/ l) alkali or acid against pH values as ordinates. Byreference to the titration curve chart, a person Without any chemicaltraining can, in a moment, add to cc. of the universal buffer solutionthe necessary Volume of N/5 acid or alkali to make not only the usualpI-I standards differing by pI- values of 0.2 but any other solutions ofany desired pH value Within 0.01, representing an unusually high orderof accuracy. By making these solutions up to a constant volume, such as20 cc. and adding an appropriate standard indicator tablet or otherstandard quantities of desired indicators, an excellent series ofcolorimetric standards is obtained.l Any desired standard for a hydrogenelectrode is also readily obtained. Formerly, in order to cover thedesired range of pH values, for example from 2 to 10 in small increments(say 0.2 pI-I) it was necessary to take 10 cc. portions of a number(usually ve or six) of buffer solutions, each of which covered a narrowpH range, and add successively increasing amounts ofalkali (or acid).The result of the tremendous amount of labor and painstaking andaccurate work Was some forty solutions which quickly became uselessbecause, being' unsterile, they Were Soon (in from two to four Weeks)rendered useless due to bacterial contamination which modified theirpI-I values. A great advance in the art is thus made by replacing forty,deteriorating solutions by a single, reliable, sterile, indefinitelystable universal buffer solution of any desired pH value.

Many solutions have been made, tested and used for these properties, andthe following are given as exemplary. A mixture of pyrophosphates,acetates, and iorrniates of sodium containing irom l-5% of formaldehydewas formed to remain sterile, but the pI-I value varied someachat dueprobably to hydration of the pyrophosphate. A solution of N/lO disodiumphosphate, iii/20 sodium acetate, N /20 sodium formate, N/ 10 sodiumsulphocarbolate, and N/20O thymol formed an excellent universal buffersolution which remained colorless, sterile and constant in pl-I valuebut its pH 'titration curve was not quite as near a straight line as wasdesired. An excellent mixture for most purposes was found to be thatresulting from the solution of 110 grams of sodium formate, 344 gramsor" sodium phenolsulphonate, 663 grains of potassium dihydrogenphosphate, 440 grams oi potassium acetate, and 30 grams of thymol in 44liters of the ypurest water. The appended curve shown in Fig. 6, SerialNo. 624,277 and Fig. 3 of the present case, is for such a solution.

Such a universal solution can be used in conimercial operations or inanalytical chemistry. in the latter case, by way of example, one of theconstituents can be a phosphate, oxalate, sulphide, carbonate, etc.,which Will give desired precipitates with added metallic radicles at theproper pH value. Or, furthermore, the universal buiier solution maycontain such metallic radicles combined with organic acids, singly or inmixture, the salts of these acids being soluble and suitable for givingany desired pi-i values; phosphates, carbonates, oxalates, sulphides,etc., can then be added to the universal buffer solution regulated toany pH value, in order to study and control the pI-I range ofprecipitation of such phosphates, hydroxides, etc. Similarly, the eiectoi pli values on oxidation-reduction reactions can be studied and usedfor their control, the oxidation-reduction agents themselves beingbuffers or not, as desired.

In the development oi overlapping buiiers covering the most usefulpI-Iranges, pI-I2-10, especially the alkaline ranges from pI-I'Z to 10,I have found that the dearth of suitable acids can be augmentedadmirably by the use of l-phenol-/i-sulphonic acid and the correspondingl-phenol-2- chlor-4-sulphonic acid, 1-phenol-2,6-dichlor-4- sulphonicacid and l-phenol-2,4-dichlo;'6-sulphonic acid. Other similar compoundscan be used in which the halogen or other negative groups cause anincrease in the ionization constant and the lower the pH range of thephenolic group. Similarly the introduction of positive alkyl or aminogroups ortho or para to phenolic groups decreases the ionizationconstant and raises the pI-I range. The introduction of one chlorine orbromine atom ortho to the hydroxyl group increases the ionizationconstant from about 102 to 162-5 fold or lowers the mid points of theuseful pH ranges from about 1.8 to 2.5 units. Two chlorine or bromineatoms ortho to the phenolic group lower the mid points of the useful pHranges about 4 pH units. For example, the curve in the annexed Fig. 8shows that in titrating 20 cc. of M/20 sodium 2-chlorphenol-4-sulphonate With 5 co. of N/5 alkali the useful pI-I range extends fromabout 6 to 8 or about 2 pH units lower than the pH range 8-10 covered bysodium phenol-4-sulphonate. Not only are these halogenated phenolderivatives very soluble but they are excellent for killing bacteria,lice, ticks, etc., and solutions thereof can be adjusted at any desiredpH value for each phenolic derivative. The phenol group itself in eachcase covers a denite pI-I range characteristic of itself, and the pHregions covered by the sharp inflections between the sulphonic acidgroup and the phenolic group can be smoothed out and regulated by theaddition to such phenolic substances of another buffer, in sufcientamount such as quantities equivalent to the phenolic group, whosecharacteristic pH range overlaps the sharp inflection pI-I range of thephenolic group. This principle is used in making my universal buffersolutions and tablets as set forth in Ser. No. 624,277.

The water used for making the dilutions or solutions of any of thesebuffers or indicators should be substantially free from even traces ofcarbon dioxide and buffer or neutral salts and hence should be pH watersuch as I prepare in my superpure water still. Unless the carbonic acidand carbonate-bicarbonic salts and other buffer materials present inordinary tap water are absent in the distillate therefrom thesechemicals change the pH of the phthalate and other buffer solutions aswell as of the indicators and buerindicator mixtures diluted with suchdistilled water. The amount of the pI-I change depends upon the relativeconcentrations and ionization constants of these buffers and indicatorsand upon the pH and degree of neutralization of each.

Ordinary puried city water as well as neutral river waters contain freecarbonic acid, carbonates, bicarbonates, lime salts, alkali metal salts,sulphates, chlorides, organic matter, sometimes free chlorine and otherconstituents all of which vary with the character of the country fromwhich the water iiows. Many of these chemicals are buffer materials inthat they change the hydrogen icn concentration of acids, bases, acidsalts, and basic salts. Fig. 4 shows two curves illustrating the effectof carbonic acid on tap water and on my superpure water and shows theirdifference in pI-I lwhich is also equal to in which VH is the number ofliters of solvent containingr one gram equivalent of hydrogen ions CH isthe concentration of the hydrogen ions and logro is the usual expressionfor the logarithm on the base i0. The lower curve A expresses therelation between the pH (as ordinates) of a certain tap water saturatedwith carbon dioxide and the dilution thereof (as abscissae) by the sametap water free from free dioxide by passage therethrough of CO2- freepure air until a constant pI-I of about 9.6 in this case is obtained. Itis seen that this particular tap water has a pH value of 4.7 whensaturated r with CO2 and that in the dilution the pH increases in nearlya straight line until a pH of 6.9 or 7.0 is obtained, which is alsoabout the pH of this tap water when collected from the faucet. As thissolution is diluted still further with CO2- free tap water there is asharp break upward in the pI-I until pH 9 or thereabouts is reachedafter which the pH curves slowly upward until a pi-I of about 9.6 isreached at a dilution of .1.:500 to 1:1000 after which it remainsconstant at pI-i' 9.6 on further dilution. It is noted that pI-l 8.1 is

obtained when this tap water is in equilibrium with ordinary air and thecarbon dioxide therein whereas superpure water gives a pH of 5.7 underthe same conditions. Turning now to the upper curve B it is noted thatsuperpure water gives pI-l 3.7 when Saturated with carbon dioxide, givespH 5.7 when in equilibrium with ordinary air and the carbon dioxidetherein and gives pH 7.0 when thoroughly scrubbed with air freed fromCO2 by passage through a soda lime scrubber. The curve B (Fig. 4) showsall the pH values intermediate between 3.7 and 7.0 given by diluting aunit volurne of superpure water saturated with carbon dioxide with theunit volumes of superpure CO2- free water ywhose logia values shown asabscissae. It is seen that curve B has no sharp inflections shown incurve A because4 the superpure water has no basic buffer materials foundin the tap water and of course partly neutralizing the carbonic acid togive the sharp inflection in curve A (Fig. 4) between about pH 7.0 and9.0. The absence of this inflection in curve B proves beyond questionthat my method of preparation of the superpure water in my superpurestill removes up to the last traces the buffer salts and carbonic acidpresent in the tap water used as a source of the superpure distilledwater. This fact is further shown by reference hereinbelow to curves inFigs. 5, 6 and 7, giving the relation of the pI-I value to the dilutionwhen phosphate and phthalate buffers or a culture medium such as Fermissolution is diluted with ordinary distilled water having a pH of about 5from the CO2 and sprayed over salts therein or with pH 5.7 waterobtained from my superpure water brought into equilibrium with ordinaryair or with my superpure pI-I 7.0 water: the curves show clearly thatthe addition of solutions of indicators (adjusted 0.2 pH apart by myisohydric indicator method described above and illustrated in thecurves) to these diluted buffer solutions indi- Cates Widely differentpH values depending upon whether ordinary distilled pH 5.0 water, airequilibrium pH 5.7 water or superpure pI-I 7.0 water is used in thedilutions and also upon the addition of indicator having the same pl-Ias the buffer solutions by my ischydric indicator method to get thecorrect pH of the buifer solution or buffer-indicator solution undertest, see Fig. 5. The curves show further that the superpure distilledwater is the only kind that will give correct results when I dilutebiological or industrial or other liquors. The use of ordinary distilledpH 5.0 water or even of air equilibrium pH 5.7 water, heretofore calledhigh grade conductivity water, begins to produce considerable errors atdilutions of 1: 10 to 1:100 if the pli-I of the buffer solution is aboveabout pH 5.0. It is only when the buer solution is quite acid such asabout pH $5.75 illustrated in Fig. 6 that no appreciable error is causedby the CO2 in even ordinary distilled pH 5.0 water whereas considerablediierences in pI-l readings are caused by the adjustment of the bromphenol blue indicator over its useful range of about pH 3.4-4.6 inintervals of 0.2 pH. This last fact shows the necessity in ne work ofadjusting the pH of the indicator solution itself and using it in thisisohydric indicator technique and also of using only superpure water ofapproximately pH 7.0 in all work on buffer solutions. This appliesin'all kinds of scientific and technical problems involving solutionsfor ore iiotations, the manufacture of pure beverages, solutions ofsugar during the reiining processes and the tests of the small amountsof the buffer impurities remaining in such refined sugar, and tests onboiler waters, sewage, factory waste liquors, natural and purifieddrinking water, plant extracts, culture media like corn meal agar orother weakly buffered media and countless other scientific and indusmtrial liquids. I-'n testing all such dilute and weakly bufferedsolutions colorimetrically the indicator must be ischydric with thesolution to which it is added, or have the same pH as the latter, togive the correct colorimetric pH value.

The curves on Fig. appended hereto illustrate how much purer and betterthe pH 7 water obtained in my superpure still is than air equilibrium ofpH 5.7 water and how much better than pH 5.0 water. These curves showdistinctly that pH 5 water has a great deal of carbon dioxide and givesvery erroneous pH results in measurements of phosphate solutions dilutedfor comparative purposes, with pH 5 water, pH 5.7 water and 7 water.This also shows that the indicator solutions themselves must be adjustedand that entirely7 diierent readings are obtained with differentadjustments of the same indicator; all three sets of curves showdifferences due to this adjustment of the indicator. With the use of thesuperpure distilled water having a pH 7.0 and an indicator having thesame pH as the buffer with which it is mixed the correct reading can beobtained. The use of pI-I 5.0 water or air equilibrium pH 5.7 watergives entirely erroneous results whether the indicator is adjusted ornot.

Fig. 1 illustrates the change in pH of the second acid group ofphosphoric acid shown as ordinate, with change in mol fractions ofalkali added to neutralize that acid group.

Fig. 2 shows how the addition. of two moles of alkali to phthalic acidbegins to form sodium acid phthalate, curve B, as the free phthalicacid, distance A, is used up. Finally the formation of disodiumphthalate begins, curve C, and the sodium acid salt then decreases inconcentration. The total curve BDC gives the relation of pH to molesalkali.

Fig. 3 shows how the pI-I of 10 cc. of universal buffer solution riseswhen treated with the indicated cc. of standard alkali and diluted to 20cc., but decreases when treated similarly with standard acid.

Fig. 4 shows how the 3.7 pH value of water saturated with CO2 increaseswith dilution of the CO2 with pure pH 7.00 water, and that similarchanges take place when pI-I 4.7 tap water saturated with CO2 is dilutedwith the same tap water freed of CO2. The abscissae 1-2, 1-5, etc.,indicate that one volume of the CO2-saturated water is diluted two-fold,iive-fold, etc.

Fig. 5 shows that a molar discdium phosphate solution having a pI-Iabout 7 .8 does not change appreciably in pH on dilution with pH 5.0,5.7 or 7.0 water, but further dilution shows distinct zones of pH withthe appropriate indicator such as phenol red adjusted in 0.2 pI-I stepsfrom its lowest to highest useful pH values.

Fig'. 6 shows that a molar potassium acid phthalate solution having a pHvalue about 3.75 is not changed materially by teniold dilution with pH5.0 water and use oi bromphenol blue, pH 3.4 to 4.6 the indicatorcovering this range, adjusted in 0.2 pH steps. After i0 fold dilutionthe elfect of the pH adjustment of the indicator is shown in thecorresponding divergence of the pH when compared against pH standards,the pH 5.0 water gradually raising the pH.

Fig. 7 shows that the electrometric pI-I of the phosphate rises ondilution, and crosses the colorimetric pH values obtained with 0.05buiercolor standards. The 0.001 molar buffer-color standards would givethe true colorimetric pH value of the diluted phosphate.

Fig. 8 shows the relation of the pH of 2-chlorphenol-4-sulphonic acid toits degree of neutralization in going from free acid to the sulphonatesalt (up to pH 4) and then the phenolate salt (up to pH 9) The termstabilized hydrogen ion standard or buer standard is intended to embracesolutions or solids used for analytical, biological, medical, commercialor other purposes.

An indicator as herein used is a dye or colored compound with acid orbasic buier properties. In the claims the term buffer includes bothcolorless and color-bearing buiers.

What I claim is:

1. An indicator standard adjusted in pH value and in concentration whichhas a CO2 content such that it is in equilibrium with the CO2 of theatmosphere.

2. A sterile buifer standard adjusted in pH value and in concentrationwhich has a CO2 content such that it is in equilibrium with the CO2 ofthe atmosphere.

3. A new sterile buffer hydrogen ion standard adjusted in pH value andin concentration and comprising a buier disinfectant, an alkali metalcarbonate and an indicator.

4. A new buer hydrogen ion standard adjusted in pH value and inconcentration and which comprises water, a buffer, an alkali metalcarbonate, a buer disinfectant and an indicator, the latter beingpreviously adjusted in such a manner that it is isohydric with the saidbuffer.

5. The step of adding an adjusted indicator to a buffer solution whichis isohydric therewith.

6. A buffer hydrogen ion standard adjusted in pH value and inconcentration and which is substantially in equilibrium with the acidand basic buifer components of the surrounding atmosphere which affecthydrogen ion concentration.

7. A stable universal buffer adjusted as to pH value and concentrationand comprising a plurality of substances which are capable of regulatinghydrogen ion concentration in a solution, differing from one another inionization constants, so as to give overlapping useful pH ranges anduniform changes in pH value upon the addition of uniformly increasingamounts of acid or alkali and said buifer containing CO2 in an amountsuch that it is in equilibrium with the CO2 content of the surroundingatmosphere.

8. In the process of determining the pH value of an unknown solution,the steps which comprise determining colorimetrically by preliminarytests with pI-I-adjusted indicators and comparisons with pI-I colorstandards the approximate pH value of small portions of said unknownsolution, adding to another portion of the said unknown solution thepH-adjusted indicator having approximately the pH value of the unknownsolution as indicated by the above said preliminary tests, and repeatinguntil the pH value of the.l unknown solution has been determined withinany desired fraction of a pH unit.

9. A process of preparing a series of indicators standardized as to pI-Iand chosen concentration without added buffers which consists in addingto chosen concentrations of said indicators having useful colorimetricpH ranges between pI-I 0 to pH 14 the requisite amounts of standardalkali or acid and making comparisons thereof in thin layers by Beerslaw against pH color standards to form a series of indicator standardsvarying from pH 0 tc pI-I 14 by denite predetermined small increments.

10. The process of preparing a stable hydrogen ion standardsusbtantially free from salt errors which comprises diluting a buier toapproximately one-thousandth molar (0.001 M) concentration and adjustingits pI-I to the desired value, and adding thereto an indicator of chosencontr'ation and pH Value, and air-equilibriating the resultantbuffer-indicator mixture and adjusting its pH to the final desiredvalue.

SOLOMON F. AGREE.

